[H2 Chemistry] 2023 Topic 2 Chemical Bonding 2

The lecture delves into covalent bonding, contrasting it with ionic and metallic bonding, and emphasizing its complexity. It focuses on the formation of sigma and pi bonds, which result from head-on and side-on orbital overlaps respectively. The sigma bond is highlighted as the first bond to form, followed by pi bonds. The lecture also explains the prerequisites for understanding covalent bonding, which include familiarity with atomic structure and orbitals. Examples of s and p orbital overlaps are provided, such as between two hydrogen atoms for a sigma bond or a sigma bond between hydrogen and fluorine. The importance of the order of bond formation is reiterated, with sigma bonds forming before pi bonds, which is a concept not always clear in secondary school curriculums.

This section further explores the nature of sigma and pi bonds, discussing the orientation of p orbitals and their involvement in bond formation. It explains that the use of a p orbital for sigma bond formation precludes the formation of additional sigma bonds, leading to the formation of pi bonds instead. The lecture also touches on the relative strength of these bonds, with pi bonds being weaker due to less direct orbital overlap. The concept of bond energy is introduced, with an emphasis on homolytic bond breakage, where bond energy is defined as the energy required to break a mole of covalent bonds homolytically. The lecture provides examples of bond energy in hydrogen and mentions the exception of fluorine, which has a weak bond energy due to electron repulsion between lone pairs on the p orbitals.

The lecture continues with a deeper look at bond energy, explaining how it is measured and the significance of homolytic bond breakage in defining bond energy. It also discusses bond strength in relation to the extent of orbital overlap and the electrostatic hold on electrons by the nuclei. The concept of bond order is introduced as a proxy for the number of covalent bonds between atoms, and the strength of multiple bonds is compared, with double and triple bonds being stronger than single bonds. The lecture also explains the increased strength of polar covalent bonds due to electrostatic attraction and mentions that bond length and strength are inversely related, with shorter bonds being stronger.

This part of the lecture addresses exceptions to the octet rule, particularly for elements like beryllium and boron, which form electron-deficient compounds like beryllium dichloride and boron trifluoride. It also discusses the stabilization of these compounds through various means, such as polymeric forms or dimeric structures involving dative covalent bonds. The lecture then transitions into the importance of dot and cross diagrams for representing covalent bonds, especially in covalent compounds, and introduces the concept of Lewis structures.

The lecture provides a practical guide to drawing dot and cross diagrams for ionic compounds, using sodium chloride and magnesium oxide as examples. It explains the electron transfer process in ionic bonding and how to represent it in a dot and cross diagram, emphasizing the difference from secondary school teachings where valence electrons are typically shown. The lecture also covers the exercise on drawing dot and cross diagrams for various ionic compounds, highlighting the process for magnesium oxide and sodium oxide, and explaining the rationale behind the chemical formula of sodium oxide.

This section discusses exceptions to the octet rule, where molecules like beryllium dichloride and boron trifluoride have central atoms with fewer than eight electrons. It explains that these electron-deficient molecules are stabilized through different mechanisms, such as polymeric forms or dimeric structures with dative covalent bonds. The lecture also addresses the concept of octet expansion for elements in periods three and above, which can accommodate more than eight electrons due to the availability of 3d orbitals.

The lecture explores octet expansion further, providing examples of molecules like PCl5 and SF4, where central atoms have more than eight electrons due to the availability of 3d orbitals. It contrasts this with period two elements, which cannot expand their octet due to the lack of accessible 3d orbitals. The lecture also discusses molecules with unpaired electrons, such as NO2 and ClO2, where the least electronegative element tends to have the unpaired electrons, making them radicals and highly reactive.

This part of the lecture offers guidelines for drawing dot and cross diagrams, particularly for polyatomic ions like sulfate, nitrate, and ammonium. It provides a set of rules for determining the central atom, distributing charges, and ensuring that terminal atoms achieve octets. The lecture also discusses exceptions to these guidelines and emphasizes the importance of practice in drawing accurate diagrams.

The lecture introduces coordinate or dative covalent bonds, where one atom donates both electrons for a single bond, often to an atom with an empty orbital. It uses ammonium and Al2Cl6 as examples, explaining how to represent these bonds in dot and cross diagrams and structural formulas. The lecture also discusses the formation of dative covalent bonds in molecules like NO2 and ClF4-, where they help achieve stable electron configurations.

This section demonstrates how to draw dot and cross diagrams for complex ions, such as the carbonate ion and the perchlorate ion. The lecture shows the process of assigning electrons to oxygen atoms to form double bonds and distributing extra electrons from external sources. It also explains how to represent the sodium ion, which donates electrons to the carbonate ion.

The lecture concludes the discussion on dot and cross diagrams and introduces the next topic, which will cover sections six and onwards. It summarizes the importance of understanding covalent bonding and the ability to draw dot and cross diagrams as a foundation for further studies in chemistry.